I'm sure you're aware that the electrons in atoms aren't stationary, and that like charges repel while opposite charges attract? In a nutshell the van der waal's forces arise from the fact that electrons are mobile within the atom. Imagine the atom split in two halves with electrons drifting randomly between and inside them. At some points in time there are more electrons one side than on the other. When this happens the side with the most electrons becomes slightly negative in charge and the other side becomes slightly positive. The negative side repels electrons in neighboring halfs of atoms so it induces the same partial charges in neighboring atoms. The positive side also does this, but by attracting electrons in the other atoms instead of repelling them.

In hydrogen and other non-polar covalent compounds this makes the molecules behave like extremely weak magnets. The "poles" attract the opposite "poles" in neighboring molecules.

The strength of these forces is dependent on the number of electrons in the atoms which make up the compound, because atoms with more electrons have a higher probability of a charge inbalance occurring and also have larger inbalances, so for instance 3 electrons to many on one side instead of 1 will make the van der waal's force about 3 times stronger.The size of the charge inbalance is also dependent on the distance from the nucleus/distance from the edge of the atom but you wont need to know that for A level. It's also slightly different in hydrogen because the only electrons in the molecule are localised between the two hydrogen nuclei, so the fluctuations in charge tend to occur mostly in the middle of the molecule, leaving the ends slightly positive most of the time, but you also dont need this for A level. So for example, in helium there's 2 electrons and in krypton there's 36. So the forces between krypton atoms is about 18 times stronger (ignoring the effect of distance from the edge of the atom). The electrostatic force's (which is what van der waal's is) strength decreases with distance, but the ratio will always be 1:~18 in this example. By the way, if you dont need it for A level, don't put it in the answers you give ok? Well, unless its now needed (I think they changed the specs, or at least, OCR did).

Since boiling and melting points are dependent on the strength of the forces between the atoms/molecules in simple covalent and monatomic structures or in giant lattices the ion attraction's strength or the bond strengths, stronger forces/bonds/whatever will give the compound a higher melting and boiling point.

Exam answer: "The melting/boiling points increase because the compounds (atoms if monatomic) have more electrons in them, so there are greater fluctuations in partial charges and therefore stronger van der waal's forces"

They might also ask you to explain the increase in melting/boiling points of the alkanes as you go along the group in order of increasing carbon content. For this DO NOT TALK ABOUT ELECTRON AMOUNTS. Here you need to refer to the surface area of the molecules. Since the strength of the forces decreases with increasing distance, the strength is maximum when the molecules are in contact with eachother. Longer chains give a higher probability of areas coming into contact and also means more parts of the molecules can be in contact with other molecules at the same time. This increases the van der waal's strength and thereby the melting/boiling points. This applies to all compounds, but only use it to answer the questions where you're dealing with chains.

However, if they ask you why branched molecules have lower melting/boiling points than the straight chain isomers (more branches gives lower mp/bp) the answer is slightly different. Its best if you visualise packing things into a suitcase. If you have 2 equally sized suitcases, one for long straight objects and one for branched objects, and equal numbers of branched and straight objects (both are in excess) you will always be able to pack more of the straight objects than the branched objects. The same thing happens with molecules. The branched isomers take up more space and limit the surface area of contact between the two molecules, because unlike the straight chain isomers they cant slide neatly next to each other, they can only fit next to each other in a way which creates big gaps between them, and the van der waal's forces' strength decrease with distance.

If you're asked to explain the difference in mp/bp of two polar molecules don't mention van der waal's. The van der waal's forces are so tiny in comparison to the dipole attractions are hydrogen "bonds" that overall they don't have a significant effect on the strength of the attractions. It's like comparing the moves a pawn can make to those a queen can make.

For this I'm only going to use the most thermodynamically stable forms of each element in the explanation bits, because mp/bp varies with allotrope :I

The first 3 are all metals (aluminium is debateable but for this we can ignore the covalent character of it's bonds) and all of them have metallic bonding. As you know the number of delocalised electrons increases across the period (only in metals of course). So as the numbers of delocalised electrons increase so does the strength of the metallic bonding. It's also because the charges on the metal ions increases so the attractions between the ions and delocalised electrons are stronger. You may have noticed that aluminium's melting point is much lower than expected. Every period has it's own black sheep, and aluminium is period 3's one. Oddly, they all seem to appear in group 3, and group 3 has some pretty wacky chemistry - which is why they don't teach it to you at A level. It also has a higher boiling point than expected. These kind of anomalies are common to group 3, the group 3 elements are kind of rebellious like that. Silicon is the only one in the period which forms giant covalent lattices (phorphorus can in theory form fullerenes - "phosphorenes" - and aluminium has the potential to, but I don't know if it does; okay fullerenes are more like giant infinite molecules but they're the border stragglers and are irrelevant here). Anyway, the covalent bonds in silicon are really strong and there's lots of them, so it takes a lot of energy to break each one, let alone all of them. So silicon has a very high mp and bp (it's smaller than carbon's because while they can both form diamond and graphite-like structures, the carbon atom is smaller than the silicon atom so the bonding electrons are closer to the nuclei, so the bonds are stronger). It's easy to assume phosphorus forms P2 molecules like nitrogen's N2, but instead it forms tetrahedral P4 molecules (white phosphorus). These are a 3D shape, so they take up more space than a triangle does, and they're also non-polar so their only intermolecular force is van der waal's. Likewise with sulfur it's really easy to think it forms S2 molecules, but instead it forms S8 rings. These are the characteristic yellow color of sulfur (coincidentally O8 molecules - yes they do exist, but very briefly - are an orange colour) and also create the smell (in addition to sulfur oxides near them). When looked at from the side they resemble crowns. These are much narrower than P4 molecules though so they have a lot more surface area in contact with each other than P4 molecules do. They also have more electrons, so these two factors combined give it stronger van der waal's forces. Chlorine has a lower mp and bp than phosphorus because Cl2 molecules are 2 atoms smaller, and the above van der waals stuff explains why this affects the mp and bp. Finally, now that the other period 3 elements argon from the list , argon is monatomic, and by now I'm sure you know what I'm going to say.

(Original post by Peroxidation)
I'm sure you're aware that the electrons in atoms aren't stationary, and that like charges repel while opposite charges attract? In a nutshell the van der waal's forces arise from the fact that electrons are mobile within the atom. Imagine the atom split in two halves with electrons drifting randomly between and inside them. At some points in time there are more electrons one side than on the other. When this happens the side with the most electrons becomes slightly negative in charge and the other side becomes slightly positive. The negative side repels electrons in neighboring halfs of atoms so it induces the same partial charges in neighboring atoms. The positive side also does this, but by attracting electrons in the other atoms instead of repelling them.

In hydrogen and other non-polar covalent compounds this makes the molecules behave like extremely weak magnets. The "poles" attract the opposite "poles" in neighboring molecules.

The strength of these forces is dependent on the number of electrons in the atoms which make up the compound, because atoms with more electrons have a higher probability of a charge inbalance occurring and also have larger inbalances, so for instance 3 electrons to many on one side instead of 1 will make the van der waal's force about 3 times stronger.The size of the charge inbalance is also dependent on the distance from the nucleus/distance from the edge of the atom but you wont need to know that for A level. It's also slightly different in hydrogen because the only electrons in the molecule are localised between the two hydrogen nuclei, so the fluctuations in charge tend to occur mostly in the middle of the molecule, leaving the ends slightly positive most of the time, but you also dont need this for A level. So for example, in helium there's 2 electrons and in krypton there's 36. So the forces between krypton atoms is about 18 times stronger (ignoring the effect of distance from the edge of the atom). The electrostatic force's (which is what van der waal's is) strength decreases with distance, but the ratio will always be 1:~18 in this example. By the way, if you dont need it for A level, don't put it in the answers you give ok? Well, unless its now needed (I think they changed the specs, or at least, OCR did).

Since boiling and melting points are dependent on the strength of the forces between the atoms/molecules in simple covalent and monatomic structures or in giant lattices the ion attraction's strength or the bond strengths, stronger forces/bonds/whatever will give the compound a higher melting and boiling point.

Exam answer: "The melting/boiling points increase because the compounds (atoms if monatomic) have more electrons in them, so there are greater fluctuations in partial charges and therefore stronger van der waal's forces"

They might also ask you to explain the increase in melting/boiling points of the alkanes as you go along the group in order of increasing carbon content. For this DO NOT TALK ABOUT ELECTRON AMOUNTS. Here you need to refer to the surface area of the molecules. Since the strength of the forces decreases with increasing distance, the strength is maximum when the molecules are in contact with eachother. Longer chains give a higher probability of areas coming into contact and also means more parts of the molecules can be in contact with other molecules at the same time. This increases the van der waal's strength and thereby the melting/boiling points. This applies to all compounds, but only use it to answer the questions where you're dealing with chains.

However, if they ask you why branched molecules have lower melting/boiling points than the straight chain isomers (more branches gives lower mp/bp) the answer is slightly different. Its best if you visualise packing things into a suitcase. If you have 2 equally sized suitcases, one for long straight objects and one for branched objects, and equal numbers of branched and straight objects (both are in excess) you will always be able to pack more of the straight objects than the branched objects. The same thing happens with molecules. The branched isomers take up more space and limit the surface area of contact between the two molecules, because unlike the straight chain isomers they cant slide neatly next to each other, they can only fit next to each other in a way which creates big gaps between them, and the van der waal's forces' strength decrease with distance.

If you're asked to explain the difference in mp/bp of two polar molecules don't mention van der waal's. The van der waal's forces are so tiny in comparison to the dipole attractions are hydrogen "bonds" that overall they don't have a significant effect on the strength of the attractions. It's like comparing the moves a pawn can make to those a queen can make.

For this I'm only going to use the most thermodynamically stable forms of each element in the explanation bits, because mp/bp varies with allotrope :I

The first 3 are all metals (aluminium is debateable but for this we can ignore the covalent character of it's bonds) and all of them have metallic bonding. As you know the number of delocalised electrons increases across the period (only in metals of course). So as the numbers of delocalised electrons increase so does the strength of the metallic bonding. It's also because the charges on the metal ions increases so the attractions between the ions and delocalised electrons are stronger. You may have noticed that aluminium's melting point is much lower than expected. Every period has it's own black sheep, and aluminium is period 3's one. Oddly, they all seem to appear in group 3, and group 3 has some pretty wacky chemistry - which is why they don't teach it to you at A level. It also has a higher boiling point than expected. These kind of anomalies are common to group 3, the group 3 elements are kind of rebellious like that. Silicon is the only one in the period which forms giant covalent lattices (phorphorus can in theory form fullerenes - "phosphorenes" - and aluminium has the potential to, but I don't know if it does; okay fullerenes are more like giant infinite molecules but they're the border stragglers and are irrelevant here). Anyway, the covalent bonds in silicon are really strong and there's lots of them, so it takes a lot of energy to break each one, let alone all of them. So silicon has a very high mp and bp (it's smaller than carbon's because while they can both form diamond and graphite-like structures, the carbon atom is smaller than the silicon atom so the bonding electrons are closer to the nuclei, so the bonds are stronger). It's easy to assume phosphorus forms P2 molecules like nitrogen's N2, but instead it forms tetrahedral P4 molecules (white phosphorus). These are a 3D shape, so they take up more space than a triangle does, and they're also non-polar so their only intermolecular force is van der waal's. Likewise with sulfur it's really easy to think it forms S2 molecules, but instead it forms S8 rings. These are the characteristic yellow color of sulfur (coincidentally O8 molecules - yes they do exist, but very briefly - are an orange colour) and also create the smell (in addition to sulfur oxides near them). When looked at from the side they resemble crowns. These are much narrower than P4 molecules though so they have a lot more surface area in contact with each other than P4 molecules do. They also have more electrons, so these two factors combined give it stronger van der waal's forces. Chlorine has a lower mp and bp than phosphorus because Cl2 molecules are 2 atoms smaller, and the above van der waals stuff explains why this affects the mp and bp. Finally, now that the other period 3 elements argon from the list , argon is monatomic, and by now I'm sure you know what I'm going to say.